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Modern Atomic Theory (a.k.a. the electron chapter!) Chapter 4 1 2 Why are electrons important? • They are the “exterior” of the atom. • Although the number of protons mainly determines the number of electrons, it is the electrons and how they are arranged that determine EVERY property of a substance! • All bonding of elements into compounds is determined by the configuration (arrangement) of the atom’s electrons. • CHEMISTRY IS ALL ABOUT ELECTRONS! 3 ELECTROMAGNETIC RADIATION – CH4.1 4 • Why do we talk about electromagnetic radiation in this unit?????? • All electromagnetic radiation comes from disturbances in the electrons of atoms. • By studying E.M. radiation, we can figure out how electrons are arranged in atoms. Electromagnetic radiation. 5 6 Electromagnetic Radiation • Most subatomic particles behave as PARTICLES and obey the physics of waves. 7 Electromagnetic Radiation wavelength Visible light Amplitude wavelength Ultaviolet radiation Node 8 Electromagnetic Radiation • Waves have a frequency, per sec” or Hertz. Hz f, units are “cycles denotes wavelength, units are meters, m • All electromagnetic radiation: • f = c where c = velocity of light = 3.00 x 108 m/sec • Now, rearrange this equation for equations for f and C Do these problems showing work using G.U.E.S.S. 9 • Remember n or nano = 10-9, Speed of light = 3X108 m/s 1) Find the wavelength of a radio wave that has a frequency of 93.1megahertz – the frequency of WXRT, Chicago’s finest rock. 2) Find the frequency of a 633nm laser beam. 3) If a light beam has a wavelength of 500nm and a frquency of 6 X 1014 Hz, find its speed. The following are NOT E.M. waves, you can’t use C= 3X108m/s 4) If an ocean wave passes me at 0.1 waves every second and there is 7 m between the waves, find its speed 5) If a sound wave has a speed of 340m/s and a frequency of 200Hz, find its wavelength. 6) If a sound wave has a speed of 340m/s and a wavelength of 0.1cm, find its frequency 7) If a by shaking a rope 3 times per second a wave is made with a wavelength of 3m find its speed Electromagnetic Spectrum Long wavelength --> small frequency Short wavelength --> high frequency increasing frequency increasing wavelength 10 11 Electromagnetic Spectrum In increasing energy, ROY G BIV 12 Quantum Theory – CH 4.2 • By 1900, this wave theory of light generally accepted. But some unexplained things remained.. • Radiant energy – why does color of an object change as the temperature changes? • Why do certain elements emit only a particular color when heated or its gas has electricity passed through it? 13 Plank’s Theory • Explained radiant energy.. • “There is a fundamental restriction on the amount of energy that an object emits or absorbs” • These amounts are called “quanta” (singular is quantum) • E=hf • energy = Plank’s Constant X frequency 14 Photoelectric Effect • Einstein used Plank’s equation to explain the photoelectric effect. • Electrons are released from metal when exposed to light (used in solar cells etc.) • However, even very intense infra red light or red light did not work whereas even very faint UV light works. • Each “photon” particle of light has energy given by E=hf • Compton : Light acts as both a wave and a particle Atomic Line Emission Spectra and Niels Bohr 15 Bohr’s greatest contribution to science was in building a simple model of the atom. It was based on an understanding of the LINE EMISSION SPECTRA of excited atoms. Niels Bohr (1885-1962) Line emission spectrum 16 • When an atom is left alone, the electrons all have certain energy levels. This “rest” state is called the GROUND STATE. • When electrons receive energy from the environment (heat, electrical or kinetic), the electrons move to higher energy levels – EXCITED STATE. • It is when these excited electrons fall back to the ground state that this energy is released as light. • The bigger the fall, the higher the frequency. Making line emission spectra 17 18 Spectrum of White Light Line Emission Spectra of Excited Atoms 19 • Excited atoms emit light of only certain frequencies. The frequencies are unique to each element. • This suggests that the electrons are arranged in specific energy levels – these levels are known 20 Spectrum of Excited Hydrogen Gas 21 Line Spectra of Other Elements 22 The Electric Pickle • Excited atoms can emit light. • Here the solution in a pickle is excited electrically. The Na+ ions in the pickle juice give off light characteristic of that element. 23 Slit that allows light inside Light Spectrum Lab! Line up the slit so that it is parallel with the spectrum tube (light bulb) Scale Light Spectrum Lab! • Run electricity through various gases, creating light • Look at the light using a spectroscope to separate the light into its component colors • Using colored pencils, draw the line spectra (all of the lines) and determine the wavelength of the three brightest lines • Once you line up the slit with the light, then look to the scale on the right. You should see the colored lines under the scale. Slit that allows light inside Eyepiece 24 Scale Light Spectrum Lab! 25 26 Atomic Spectra One view of atomic structure in early 20th century was that an electron (e-) traveled about the nucleus in an orbit. 27 Atomic Spectra and Bohr Bohr said classical view is wrong. Need a new theory — now called QUANTUM or WAVE MECHANICS. e- can only exist in certain discrete orbits e- is restricted to QUANTIZED energy state (quanta = bundles of energy) Quantum or Wave Mechanics 28 Schrodinger applied idea of ebehaving as a wave to the problem of electrons in atoms. He developed the WAVE EQUATION Solution gives set of math expressions called WAVE E. Schrodinger FUNCTIONS, 1887-1961 Each describes an allowed energy state of an e- Heisenberg Uncertainty Principle W. Heisenberg 1901-1976 Problem of defining nature of electrons in atoms solved by W. Heisenberg. Cannot simultaneously define the position and momentum (= m•v) of an electron. We define e- energy exactly but accept limitation that we do not know exact position. 29 Arrangement of Electrons in Atoms Electrons in atoms are arranged as LEVELS (n) SUBLEVELS (l) ORBITALS (ml) 30 QUANTUM NUMBERS The shape, size, and energy of each orbital is a function of 3 quantum numbers which describe the location of an electron within an atom or ion n (principal) ---> energy level l (orbital) ---> shape of orbital ml (magnetic) ---> designates a particular suborbital The fourth quantum number is not derived from the wave function s (spin) ---> spin of the electron (clockwise or counterclockwise: ½ or – ½) 31 QUANTUM NUMBERS So… if two electrons are in the same place at the same time, they must be repelling, so at least the spin quantum number is different! The Pauli Exclusion Principle says that no two electrons within an atom (or ion) can have the same four quantum numbers. If two electrons are in the same energy level, the same sublevel, and the same orbital, they must repel. Think of the 4 quantum numbers as the address of an electron… Country > State > City > Street 32 33 Energy Levels • Each energy level has a number called the PRINCIPAL QUANTUM NUMBER, n • Currently n can be 1 thru 7, because there are 7 periods on the periodic table 34 Energy Levels n=1 n=2 n=3 n=4 Relative sizes of the spherical 1s, 2s, and 3s orbitals of hydrogen. 35 36 Types of Orbitals • The most probable area to find these electrons takes on a shape • So far, we have 4 shapes. They are named s, p, d, and f. • No more than 2 e- assigned to an orbital – one spins clockwise, one spins counterclockwise Types of Orbitals (l) s orbital p orbital d orbital 37 p Orbitals 38 this is a p sublevel with 3 orbitals These are called x, y, and z 3py orbital There is a PLANAR NODE thru the nucleus, which is an area of zero probability of finding an electron p Orbitals • The three p orbitals lie 90o apart in space 39 d Orbitals • d sublevel has 5 orbitals 40 41 The shapes and labels of the five 3d orbitals. 42 f Orbitals For l = 3, ---> f sublevel with 7 orbitals Diagonal Rule • Must be able to write it for the test! This will be question #1 ! Without it, you will not get correct answers ! • The diagonal rule is a memory device that helps you remember the order of the filling of the orbitals from lowest energy to highest energy • Aufbau Principle states that electrons fill from the lowest possible energy to the highest energy 43 Diagonal Rule – showing sublevels 44 Steps: 1s 2s 3s 1. Write the energy levels top to bottom. 2. Write the orbitals in s, p, d, f order. Write the same number of orbitals as the energy level. 3. Draw diagonal lines from the top right to the bottom left. 4. To get the correct order, 2p 3p 3d follow the arrows! 4s 4p 4d 4f 5s 5p 5d 5f 5g? 6s 6p 6d 6f 6g? 6h? 7s 7p 7d 7f 7g? 7h? By this point, we are past the current periodic table so we can stop. 7i? Rules of electron orbital filling 45 • Aufbau Principle states that electrons fill from the lowest possible energy to the highest energy. • Hund’s Rule In orbitals of EQUAL ENERGY (p, d, and f), place one electron in each orbital before making any pairs. All single electrons must spin the same way • Pauli Exclusion Principle says that no two electrons within an atom (or ion) can have the same four quantum numbers. • The quantum #’s refer to the principle energy level (n=1->6), sublevel (spdf), orbital and spin (up or down) 46 Why are d and f orbitals always in lower energy levels? • d and f orbitals require LARGE amounts of energy • It’s better (lower in energy) to skip a sublevel that requires a large amount of energy (d and f orbtials) for one in a higher level but lower energy This is the reason for the diagonal rule! BE SURE TO FOLLOW THE ARROWS IN ORDER! How many electrons can be in a sublevel? 47 Remember: A maximum of two electrons can be placed in an orbital. s orbitals p orbitals d orbitals f orbitals Number of orbitals 1 3 5 Number of electrons 2 6 10 7 14 Electron Configurations A list of all the electrons in an atom (or ion) • Must go in order (Aufbau principle) • 2 electrons per orbital, maximum • We need electron configurations so that we can determine the number of electrons in the outermost energy level. These are called valence electrons. • The number of valence electrons determines how many and what this atom (or ion) can bond to in order to make a molecule 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14… etc. 48 Electron Configurations 4 2p Energy Level Number of electrons in the sublevel Sublevel 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14… etc. 49 Let’s Try It! • Write the electron configuration for the following elements: H Li N Ne K Zn Pb 50 51 An excited lithium atom emitting a photon of red light to drop to a lower energy state. An excited H atom returns to a lower energy level. 52 Orbitals and the Periodic Table • Orbitals grouped in s, p, d, and f orbitals (sharp, proximal, diffuse, and fundamental) s orbitals f orbitals d orbitals p orbitals 53 Shorthand Notation • A way of abbreviating long electron configurations • Since we are only concerned about the outermost electrons, we can skip to places we know are completely full (noble gases), and then finish the configuration 54 Shorthand Notation • Step 1: It’s the Showcase Showdown! Find the closest noble gas to the atom (or ion), WITHOUT GOING OVER the number of electrons in the atom (or ion). Write the noble gas in brackets [ ]. • Step 2: Find where to resume by finding the next energy level. • Step 3: Resume the configuration until it’s finished. 55 Shorthand Notation • Chlorine – Longhand is 1s2 2s2 2p6 3s2 3p5 You can abbreviate the first 10 electrons with a noble gas, Neon. [Ne] replaces 1s2 2s2 2p6 The next energy level after Neon is 3 So you start at level 3 on the diagonal rule (all levels start with s) and finish the configuration by adding 7 more electrons to bring the total to 17 [Ne] 3s2 3p5 56 57 Practice Shorthand Notation • Write the shorthand notation for each of the following atoms: Cl K Ca I Bi Hg Ag Al Valence Electrons Electrons are divided between core and valence electrons B 1s2 2s2 2p1 Core = [He] , valence = 2s2 2p1 Br [Ar] 3d10 4s2 4p5 Core = [Ar] 3d10 , valence = 4s2 4p5 58 Rules of the Game No. of valence electrons of a main group atom = Group number (for A groups) Atoms like to either empty or fill their outermost level. Since the outer level contains two s electrons and six p electrons (d & f are always in lower levels), the optimum number of electrons is eight. This is called the octet rule. 59 60 Keep an Eye On Those Ions! • Electrons are lost or gained like they always are with ions… negative ions have gained electrons, positive ions have lost electrons • The electrons that are lost or gained should be added/removed from the highest energy level (not the highest orbital in energy!) 61 Keep an Eye On Those Ions! • Tin Atom: [Kr] 5s2 4d10 5p2 Sn+4 ion: [Kr] 4d10 Sn+2 ion: [Kr] 5s2 4d10 Note that the electrons came out of the highest energy level, not the highest energy orbital! 62 Keep an Eye On Those Ions! • Bromine Atom: [Ar] 4s2 3d10 4p5 Br- ion: [Ar] 4s2 3d10 4p6 Note that the electrons went into the highest energy level, not the highest energy orbital! Try Some Ions! • Write the longhand notation for these: FLi+ Mg+2 • Write the shorthand notation for these: BrBa+2 Al+3 63 (HONORS only) Exceptions to the Aufbau Principle • Remember d and f orbitals require LARGE amounts of energy • If we can’t fill these sublevels, then the next best thing is to be HALF full (one electron in each orbital in the sublevel) • There are many exceptions, but the most common ones are d4 and d9 For the purposes of this class, we are going to assume that ALL atoms (or ions) that end in d4 or d9 are exceptions to the rule. This may or may not be true, it just depends on the atom. 64 (HONORS only) Exceptions to the Aufbau Principle 65 d4 is one electron short of being HALF full In order to become more stable (require less energy), one of the closest s electrons will actually go into the d, making it d5 instead of d4. For example: Cr would be [Ar] 4s2 3d4, but since this ends exactly with a d4 it is an exception to the rule. Thus, Cr should be [Ar] 4s1 3d5. Procedure: Find the closest s orbital. Steal one electron from it, and add it to the d. (HONORS only) 66 Exceptions to the Aufbau Principle OK, so this helps the d, but what about the poor s orbital that loses an electron? Remember, half full is good… and when an s loses 1, it too becomes half full! So… having the s half full and the d half full is usually lower in energy than having the s full and the d to have one empty orbital. (HONORS only) Exceptions to the Aufbau Principle 67 d9 is one electron short of being full Just like d4, one of the closest s electrons will go into the d, this time making it d10 instead of d9. For example: Au would be [Xe] 6s2 4f14 5d9, but since this ends exactly with a d9 it is an exception to the rule. Thus, Au should be [Xe] 6s1 4f14 5d10. Procedure: Same as before! Find the closest s orbital. Steal one electron from it, and add it to the d. (HONORS only) 68 Try These! • Write the shorthand notation for: Cu W Au Orbital Diagrams • Graphical representation of an electron configuration • One arrow represents one electron • Shows spin and which orbital within a sublevel • Same rules as before (Aufbau principle, d4 and d9 exceptions, two electrons in each orbital, etc. etc.) 69 Orbital Diagrams • One additional rule: Hund’s Rule – In orbitals of EQUAL ENERGY (p, d, and f), place one electron in each orbital before making any pairs – All single electrons must spin the same way • I nickname this rule the “Monopoly Rule” • In Monopoly, you have to build houses EVENLY. You can not put 2 houses on a property until all the properties has at least 1 house. 70 Lithium 71 Group 1A Atomic number = 3 1s22s1 ---> 3 total electrons 3p 3s 2p 2s 1s Carbon 3p 3s 2p 2s 1s Group 4A Atomic number = 6 1s2 2s2 2p2 ---> 6 total electrons Here we see for the first time HUND’S RULE. When placing electrons in a set of orbitals having the same energy, we place them singly as long as possible. 72 Lanthanide Element Configurations 4f orbitals used for Ce - Lu and 5f for Th - Lr 73 74 Draw these orbital diagrams! • Oxygen (O) • Chromium (Cr) • Mercury (Hg) Ion Configurations To form anions from elements, add 1 or more e- from the highest sublevel. P [Ne] 3s2 3p3 + 3e- ---> P3- [Ne] 3s2 3p6 or [Ar] 3p 3p 3s 3s 2p 2p 2s 2s 1s 1s 75 General Periodic Trends • Atomic and ionic size • Ionization energy • Electronegativity Higher effective nuclear charge Electrons held more tightly Larger orbitals. Electrons held less tightly. 76 77 Atomic Size • Size goes UP on going down a group. • Because electrons are added further from the nucleus, there is less attraction. This is due to additional energy levels and the shielding effect. Each additional energy level “shields” the electrons from being pulled in toward the nucleus. • Size goes DOWN on going across a period. 78 Atomic Size Size decreases across a period owing to increase in the positive charge from the protons. Each added electron feels a greater and greater + charge because the protons are pulling in the same direction, where the electrons are scattered. Large Small 79 Which is Bigger? • Na or K ? • Na or Mg ? • Al or I ? Ion Sizes Li,152 pm 3e and 3p Does+ the size go up+ or down Li , 60 pm when an 2e and 3losing p electron to form a cation? 80 81 Ion Sizes + Li,152 pm 3e and 3p Li + , 78 pm 2e and 3 p Forming a cation. • CATIONS are SMALLER than the atoms from which they come. • The electron/proton attraction has gone UP and so size DECREASES. Ion Sizes Does the size go up or down when gaining an electron to form an anion? 82 83 Ion Sizes F, 71 pm 9e and 9p F- , 133 pm 10 e and 9 p Forming an anion. • ANIONS are LARGER than the atoms from which they come. • The electron/proton attraction has gone DOWN and so size INCREASES. • Trends in ion sizes are the same as atom sizes. Trends in Ion Sizes Figure 8.13 84 85 Which is Bigger? • Cl or Cl- ? • K+ or K ? • Ca or Ca+2 ? • I- or Br- ? Ionization Energy 86 IE = energy required to remove an electron from an atom (in the gas phase). Mg (g) + 738 kJ ---> Mg+ (g) + e- This is called the FIRST ionization energy because we removed only the OUTERMOST electron Mg+ (g) + 1451 kJ ---> Mg2+ (g) + eThis is the SECOND IE. Trends in Ionization Energy • IE increases across a period because the positive charge increases. • Metals lose electrons more easily than nonmetals. • Nonmetals lose electrons with difficulty (they like to GAIN electrons). 87 88 Trends in Ionization Energy • IE increases UP a group • Because size increases (Shielding Effect) 89 Which has a higher 1st ionization energy? • Mg or Ca ? • Al or S ? • Cs or Ba ? 90 Electronegativity, is a measure of the ability of an atom in a molecule to attract electrons to itself. Concept proposed by Linus Pauling 1901-1994 Periodic Trends: Electronegativity • In a group: Atoms with fewer energy levels can attract electrons better (less shielding). So, electronegativity increases UP a group of elements. • In a period: More protons, while the energy levels are the same, means atoms can better attract electrons. So, electronegativity increases RIGHT in a period of elements. 91 Electronegativity 92 93 Which is more electronegative? • F or Cl ? • Na or K ? • Sn or I ? 94 The End !!!!!!!!!!!!!!!!!!!